Organic Chemistry Made Easy

“Organic Chemistry Made Easy” attempts to help students throughout the world acquire a comprehensive understanding of important concepts in organic chemistry. Students will be able to predict products in a chemical reaction (including the stereochemistry of the products), and students will be able to interpret the spectra of organic compounds.
The intent of this blog is to provide a comprehensive course in elementary organic chemistry.

Wednesday, December 30, 2009

Organic Chemistry is the study of compounds containing carbon atoms. Inorganic compounds are obtained from minerals, and up to about 1850, scientists believed that organic compounds came from living substances or organisms.

The one thing that all organic compounds have in common is the element carbon. The science of organic chemistry didn’t evolve until around the first quarter of the nineteenth century, and organic chemistry development from that point can be divided into three overlapping periods. The first period can be considered as the general advancement of organic chemistry where chemical transformations are reproducible and proceed independent of supernatural influences; therefore, leading to the abandonment of the “vital force” theory. The abandonment of “vitalistic theory” was initiated in 1828 when Frederick Wöhler performed the following experiment:

Wöhler’s experiment was somewhat obscure, because he obtained the ammonium cyanate from the calcinations of bones; therefore, he wasn’t positive that his experiment was not contaminated. He reported his work in an unbelieving circumspect language. Consequently, it took several years before the “vital force” theory was truly dead.By 1850, the “vital force” theory was dead. This was primarily due to Kolbe’s experiment where he converted chloroacetic acid, compound I, to acetic acid, compound II:

The third period of organic chemistry came in 1859 with the birth of structural theory. Berzelius, Wöhler’s mentor, and Wöhler were curious about the conversion of ammonium cyanate to urea from a structural perspective. Structural theory was independently proposed by August Kekulé, from Germany; Archibald Couper, from Scotland; and Alexander Butlerov, from Russia.Using structural information learned in general chemistry; the following is a more detailed representation of Wöhler’s experiment:

With this transformation, it is clear that a great deal of changes must be occurring.For example, the nature of the hybridization about the carbon atom in the cyanate changes from 2sp to 2sp2. Also the hybridization about both nitrogen atoms in the cyanate change. The nitrogen of the ammonium ion changes from 2sp3 to 2sp2 , and the nitrogen in the cyanate changes from 2sp to 2sp2. In addition the bond angles also under transformation. These differences will be examined momentarily.

The third period of the development of organic chemistry involves the use of modern instrumental techniques for separating, analyzing, and identifying organic compounds. A considerable amount of time will be devoted to some of these techniques, e.g., mass spectrometry, infrared spectrophotometry, nuclear magnetic resonance spectrometry, and UV-visible spectrophotmetry.

Most organic compounds are synthesized from a variety of inorganic substance, e.g., carbonates or cyanides and primarily from other organic compounds. Petroleum and coal are two large reservoirs for obtaining organic compounds. There are more than a million organic compounds known compared to just thousands of inorganic compounds. The organic compounds outnumber inorganic compounds, because of the virtually unique ability of the carbon atom to expand its octet of electrons.

Let’s examine structural theory in more detail. Structural theory takes into consideration the number of atoms comprising the molecule, the order of attachment of the atoms, the electrons holding the atoms together, and the shapes and sizes of the molecules.

Before 1926, the important aspects of the nature of the chemical bond included the ionic bond proposed by Walter Kossel in 1916 and the covalent bond described by G. N. Lewis.

Kossel and Lewis based their bonding theories on the belief that: •the positive charged nucleus is surrounded by electrons•electrons exist in shells with a designated maxima•stability is reached when the outer shell is full•ionic and covalent bonds are consequences of the ability of atoms to obtain the stability or stable electron configuration of the nearest noble gas

Lithium fluoride, LiF, forms an ionic bond where an electron on the outer shell of Li is transferred to the outer shell of fluorine. The driving force for this transfer is that the resulting lithium ion is isoelectronic with He and the resulting fluoride ion is isoelectronic with neon. These concepts are visualized in Figure 1.1.

Figure 1.1 the formation of the ionic bond

Li+ has the stable electronic configuration that is isoelectronic with He, 1s2.F- has the stable electronic configuration that is isoelectronic with Ne, 1s2 2s2 2p6.Consequently, LiF is formed by the transfer of an electron from Li to F to form LiF, an ionic compound.

Covalent bonds are formed from the sharing of electrons. For example, the hydrogen molecule is formed from the sharing an electron from one hydrogen atom with an electron from another hydrogen atom. When this sharing process occurs, both hydrogen atoms can complete their shells; therefore, reaching the stable configuration of He, 1s2. The same consideration could be given to fluorine where the outer shell electrons would reach the stable outer shell configuration of Ne, 1s2 2s2 2p6, eight(8) electrons, generally referred to as the stable octet.Figure 1.2 represents the covalent bonding in H2 and F2

Figure 1.2 covalent bond in hydrogen and fluorine

Figure 1.3 represents the covalent bonding in CF4

Figure 1.3 represents the covalent bonding in carbon tetrafluoride

In 1926, Erwin Schrödinger used operator mathematics to develop a mathematical expression that describes the motion of electrons in terms of energy equations. These equations show properties of a wave; therefore, they are referred to as wave equations. These equations cannot be solved in an exact manner, but give only approximate solutions called wave functions. The lower the energy value, the more nearly correct the wave function. The mathematical equation giving rise to the wave functions in its most simplistic form is Hψ = Eψ where H is the Hamiltonian operator, ψ is the wave function, and E is the eigenvalue or energy value. The Hamiltonian is a complicated mathematical system that operates on a wave function that describes the behavior of the electron such that when the Hamiltonian operates on the wave function, the wave function is regenerated along with an energy value. This form of mathematics is referred to as quantum mechanics and involves a type of mathematics referred to as differential equations. For example, an electron travels in one dimension a distance x’ along the x-axis. The electron behaves like a wave and its mathematical function is

where ψ is the wave function and λ is the wavelength of the electron (the electron is now behaving as a wave)Hψ = EψLet’s assume that the Hamiltonian is Therefore,

This operation is beyond the mathematical knowledge of most individuals taking organic chemistry; however, the result would be

Noticed that the wave function,, has been regenerated and the eigenfunction or energy value is

Atomic orbitals

The solution to the Schrödinger equation for an electron in a quantized orbital must have three quantum numbers in order to give rise to the appropriate eigenvalue, energy, for the electron. The three quantum numbers are the principle quantum number, n, the energy level in which the electron resides; the azimuthal or secondary quantum, l, the shape that describes the probability of finding the electron in a volume element; and the magnetic quantum number, ml, the orientation of the volume element describing the electron in reference to the Cartesian coordinate axes.

The wave equations cannot simultaneously give the exact position of the electron at any particular moment and how fast the electron is moving. Also, no precise orbital about the nucleus can be plotted by the wave equation, but they provide the probability of finding the electron within a volume element. The orbital is the region of space where an electron can be found. As indicated earlier, the shape of the orbital depends upon the energy of the electron. Consequently, the electron cloud that describes the probability of finding the electron is a blurred photograph of a rapidly moving electron. This electron cloud density is not uniform, but densest in the region where the probability of finding the electron is highest, i.e., where the average negative charge is the greatest.

Using the Schrödinger equation, the lowest energy orbital is the 1 s orbital (Figure 1.4). The next energy orbital is the 2 s orbital (Figure 1.5). After the 2s energy orbital, the next energy orbital, the 2 p, has three orientations (all energy degenerate). These degenerate energy orbitals are the 2px, the 2py and the 2pz orbitals (Figure 1.6).

Figure 1.6 the 2p orbitals

Each 2p orbital is dumbbell shaped.

Electron Configuration

Each orbital can have a maximum of two electrons, but the two electrons must have opposite spins. The spin quantum number does not result from the Schrödinger equation, but is an outcome of the Stern-Gerlach Experiment.The Stern-Gerlach Experiment showed that electrons have two spin Orientations (Figures 1.7 and 1.8).

In the Stern-Gerlach Experiment, a beam of hydrogen atoms is passed through a magnetic field. The beam of hydrogen atoms is split into two beams suggesting that the electrons in the hydrogen atoms have magnetic properties that are associated with opposite spins.

Figure 1.7 is taken from the following Website: http://library.thinkquest.org/19662/high/eng/exp-stern-gerlach.htmlFigure 1.7 a beam of hydrogen atoms splitting into two beams as it passes through a magnetic field.Figure 1.8 shows that that one beam consist of hydrogen with a spin of +½ and the other with a spin of -½

Figure 1.8 Two spins for the hydrogen electronsThis figure was taken from Ebbing and Gammon “General Chemistry (Seventh Edition)” page 313

1.The hydrogen electrons in each beam act as magnets as a consequence of their spins.2.The orientation of spins are different, i.e., one hydrogen electron spins clockwise and the other hydrogen electron spins counterclockwise.3.The spinning electrons generate small magnetic fields.4.The electron spin is susceptible to quantum restrictions, i.e., ms = +½ or ms = -½5.Two electrons can occupy a single orbital, but the electrons must have opposite spins.

To view a dynamic illustration of the Stern-Gerlach Experiment click on the following Website: http://mutuslab.cs.uwindsor.ca/schurko/nmrcourse/animations/stern-gerlach/sgpeng.htm

There are four quantum numbers, including the spin quantum number; and no two electrons can have the same set of four quantum numbers in an orbital (Pauli Exclusion Principle). If two electrons reside in the same orbital, then one electron would have a spin of +½ and the other would have a spin of -½ since n, l and ml would be the same. Also, Hund’s Rule states that one electron must be in each orbital of a set of degenerate orbitals before two electrons can occupy any of the degenerate orbitals.

Electron Configuration and Orbital Diagrams

As indicated previously, the solution to the Schrödinger equation gives the probability of finding the electron in a volume of space. These volume spaces are the s, p, d and f orbitals, and each orbital can hold only two electrons, and the two electrons must have opposite spins.

Most organic compounds will use the s and p orbitals primarily.Subshells consist of the different orbitals within a specified volume of space. For example, when n = 2, the total number of electrons that are possible in energy level 2 would be 2n2 where n equals the energy level in question. Eight electrons are possible for n = 2, i.e., 2(2)2 = 8. When n = 2, two electrons can reside in the s orbital and six electrons can reside in three degenerate p orbitals. Two subshells exist for n = 2. The s and the p subshells. Table 1.1 indicates the maximum number of electrons that can exist in each subshell.

Table 1.1 the maximum number of electrons available for each subshell *described in the section entitled “pictorial solutions to the Schrödinger equation that represent the probability of finding electrons about the nucleus of an atom”

Each orbital can hold two electrons with opposite spins.

Let’s apply this constraint to writing the electron configuration of Li. Lithium has three electrons in orbitals surrounding the nucleus.

The first energy level, n=1, can only hold two electrons [2(1)2=2]; therefore, there is only one subshell, consisting of the 1s orbital. The second energy level has two subshells, the 2s and 2p subshells. The third electron must reside in the second energy level in the lower energy 2s orbital subshell. The electron configuration for Li would be 1s22s1. Figure 1.9 is an orbital diagram that represents the potential energy relationships of the orbital electrons in the lithium atom.

Figure 1.9 the electron configuration of Li

The following arrangement of electrons in an orbital would be forbidden:

Another important rule is Hund’s rule that was proposed by Friedrich Hund. Hund’s rule states that for a set of degenerate orbitals, i.e., orbitals with the same energy, each degenerate orbital must have one electron in it before a second electron can be placed in the orbital. When a second electron is placed in the orbital, it must have an opposite spin. For example, there are three p orbitals with the same energy in the 2p subshell. If three electrons are in the 2p subshell, than there must be one electron in the 2px orbital, one electron in the 2py orbital, and one electron in the 2pz orbital before a second electron of opposite spin can be added to the set of degenerate orbitals.Nitrogen has seven electrons in orbitals surrounding the nucleus.The first energy level, n=1, can only hold two electrons [2(1)2 = 2]; therefore there is only one subshell consisting of the 1s orbital. The second energy level, n=2, can hold [2(2)2 = 8]. Since two of the seven nitrogen electrons reside in the 1s orbital, the remaining five must reside at energy level 2. Two of these electrons would be in the 2s orbital, and the remaining three would be in the higher degenerate energy 2p orbitals.

The Pauli Exclusion Principle, proposed by physicist Wolfgang Pauli, states that no two electrons can have the same set of four quantum numbers. Since two electrons within the same orbital can have the same values for n, l, and ml; therefore, the two electrons in the orbital must have opposite spins, i.e., ms=+½ and ms=-½.

The electron configuration for nitrogen is 1s22s22p1x2p1y21z

Using Hund’s rule and the Pauli Exclusion, Figure 1.10 is an orbital diagram that represents the potential energy relationships of the orbital electrons in the nitrogen atom.

Figure 1.10 electron configuration of nitrogen obeying Hund’s rule

Table 1.2 represents the configuration for hydrogen using the four quantum numbers (three from the Schrödinger equation and one from the work of Stern and Gerlach).Table 1.2 the quantum numbers for hydrogen

Table 1.3 represents the configuration for helium (He has two electrons about the nucleus). Each electron is described using the four quantum numbers (three from the Schrödinger equation and one from the work of Stern and Gerlach). Helium is a noble gas and is inert to the formation of compounds. The first energy level is complete with two electrons; therefore, making it inert to chemical reactions. electron Table 1.3: the two electrons have identical numbers for n, l, and ml, but different spin quantum numbers.

Table 1.4 represents the configuration for lithium (has three electrons about the nucleus). Each electron is described using the four quantum numbers (three from the Schrödinger equation and one from the work of Stern and Gerlach).

Table 1.4 the quantum numbers for the three lithium electronsTable 1.5 represents the configuration for beryllium (has four electrons about the nucleus). Each electron is described using the four quantum numbers (three from the Schrödinger equation and one from the work of Stern and Gerlach.Table 1.5 the quantum numbers for the four Be electrons

Table 1.6 represents the configuration for boron (has five electrons about the nucleus). Each electron is described using the four quantum numbers (three from the Schrödinger equation and one from the work of Stern and Gerlach).Table 1.6: the quantum numbers for the five B electrons and the electron configuration for B is 1s22s22p1x2p0y20z

Table 1.7 represents the configuration for carbon (C has six electrons about the nucleus). Each electron is described using the four quantum numbers (three from the Schrödinger equation and one from the work of Stern and Gerlach).electron Table 1.7: the quantum numbers for the six carbon electrons and the the electron configuration for C is 1s22s22p1x2p1y20z

Table 1.8 represents the configuration for nitrogen (has seven electrons about the nucleus). Each electron is described using the four quantum numbers (three from the Schrödinger equation and one from the work of Stern and Gerlach).electron Table 1.8: The quantum numbers for the seven electrons for N and the electron configuration for N is 1s22s22p1x2p1y21zTable 1.9 represents the configuration for oxygen (has eight electrons about the nucleus). Each electron is described using the four quantum numbers (three from the Schrödinger equation and one from the work of Stern and Gerlach).Table 1.9: The electron configuration for O is 1s22s22p2x2p1y21z

Table 1.10 represents the configuration for fluorine (has nine electrons about the nucleus). Each electron is described using the four quantum numbers (three from the Schrödinger equation and one from the work of Stern and Gerlach).Table 1.10: The quantum numbers for the nine electrons of F and the electron configuration for F is 1s22s22p2x2p2y21z

Table 1.11 represents the configuration for neon (Ne has ten electrons about the nucleus). Each electron is described using the four quantum numbers (three from the Schrödinger equation and one from the work of Stern and Gerlach). Neon is a noble gas and is inert to the formation of compounds. The second energy level is complete with eight electrons; therefore, making it inert to chemical reactions. Table 1.11: The quantum numbers for the ten electrons for Ne and the electron configuration for Ne is 1s22s22p2x2p2y22z

Aufbau Principle is the building of the electron configuration of the elements as they are identified in the periodic chart.

Molecular orbitals

Molecular orbitals result from the linear combination of atomic orbitals. This occurs when two atoms form a chemical bond. The electron pair or pairs in the molecular orbital are located in proximity to the two nuclei forming the chemical bond. The shape of the molecular orbital are related to the shapes of the atomic orbitals in the component atoms. The bonding in methane, ammonia, water and the unsaturated system are described as below:

Methane, CH4

To understand the bonding that occurs in the covalent bonds between C and H, the electron configurations of both carbon and hydrogen should be revisited. The electron configurations of carbon and hydrogen are:C 1s22s22p2H 1s1

An electron in the 2s orbital of carbon can be promoted to a 2p orbital of carbon. The promotion of an electron from the 2s atomic orbital of C to a 2p atomic orbital is illustrated in Diagram 1.1.Diagram 1.1

The 2p atomic orbital that contains the promoted electron linearly combines with the remaining two 2p atomic orbitals of carbon (each containing an electron) and the 2s atomic orbital that contains one electron to form four degenerate (equal in energy) sp3 hybridized atomic orbitals. The resulting hybridized atomic orbitals are a little lower in energy than the 2p orbitals from which they were formed and a little higher in energy than the 2s orbital from which they were formed. Each sp3 atomic orbital has one electron. Diagram 1.2 illustrates this process.Diagram 1.2

Each C-H covalent bond is formed from the linear combination of a 1s atomic orbital of H with a 2sp3 hybridized atomic orbital of carbon to form a bonding and an anti-bonding molecular orbital that describe the carbon-hydrogen bond in methane. The bonding in each C-H bond in methane is described by Diagram 1.3.Diagram 1.3

Each bond between carbon and hydrogen is a single covalent bond (referred to as a sigma, σ bond). The σ bond is formed from the linear combination of a 2sp3 hybridized atomic orbital of carbon and a 1s atomic orbital from hydrogen. In this case the single covalent bond is given the notation The result is illustrated in Diagram 1.4.

Diagram 1.4 a tetrahedral molecular with a H-C-H bond angle equal to 109.5o

The energy required to break the C-H bond is 427 kJ/mol.

Ammonia, NH3

To understand the bonding that occurs in the covalent bonds between N and H, the electron configurations of both nitrogen and hydrogen should be revisited. The electron configurations of nitrogen and hydrogen are:N 1s22s22p3H 1s1

In the case of nitrogen, there is no need to promote an electron in the 2s orbital to a 2p orbital. The only energy requirement would be the need to hybridize the three p orbitals, each containing one electron, and the lower energy 2s orbital containing two electrons to form four degenerate 2sp3 hybridized orbitals as illustrated in Diagram 1.5.Diagram 1.5

The four 2sp3 hybridized orbitals are a little lower in energy than the 2p orbitals from which they were formed and a little higher in energy than the 2s orbital from which they were formed. Three 2sp3 atomic orbitals have one electron each and one 2sp3 atomic orbital has two electrons.Three N-H covalent bonds are formed from the linear combination a 1s atomic orbital of H with a 2sp3 hybridized atomic orbital of nitrogen to form a bonding molecular orbital and an anti-bonding molecular orbital that describe the nitrogen-hydrogen bond in ammonia. The bonding in each N-H bond in ammonia is described by Diagram 1.6.Diagram 1.6

Each bond between nitrogen and hydrogen is a single covalent bond (referred to as a sigma, σ bond). The σ bond is formed from the linear combination of a 2sp3 hybridized atomic orbital of nitrogen and a 1s atomic orbital from hydrogen. In this case the single covalent bond is given the notationDiagram 1.7 a trigonal pyramidal molecule with an H-N-H bond angle approximately 107o

Water, H2O

To understand the bonding that occurs in the covalent bonds between O and H, the electron configurations of both oxygen and hydrogen should be revisited:O 1s22s22p4H 1s1

In the case of oxygen, there is no need to promote an electron in the 2s orbital to a 2p orbital. The only energy requirement would be the need to hybridize the three p orbitals two of which contain one electron each and one containing two electrons and the lower energy 2s orbital containing two electrons to form four degenerate 2sp3 hybridized orbitals as illustrated in Diagram 1.8.Diagram 1.8

The four 2sp3 hybridized orbitals are a little lower in energy than the 2p orbitals from which they were formed and a little higher in energy than the 2s orbital from which they were formed. Two of the 2sp3 atomic orbitals have one electron each and two of the 2sp3 atomic orbitals have two electrons each.

The two O-H covalent bonds are formed from the linear combination of a 1s atomic orbital of H with a 2sp3 hybridized atomic orbital of oxygen to form a bonding molecular orbital and an anti-bonding molecular orbital that describe the oxygen-hydrogen bond in water. The bonding in each O-H bond in water is described by Diagram 1.9.Diagram 1.9

Each bond between oxygen and hydrogen is a single covalent bond (referred to as a sigma, σ bond). The σ bond is formed from the linear combination of a 2sp3,/sup> hybridized atomic orbital of oxygen and a 1s atomic orbital from hydrogen. In this case the single covalent bond is given the notation Diagram 1.10 a bent molecule with a H-O-H bond angle equal to 105o

In order to understand the nature of the double bond, let’s revisit the electron configuration of carbon. This configuration is 1s22s22p2. An electron in the 2s orbital of carbon can be promoted to the 2p orbital. Diagram 1.11 is an illustration of this process.The 2p atomic orbital that contains the promoted electron linearly combines with the 2s atomic orbital that contains one electron to form three degenerate 2sp2 hybridized atomic orbitals. Diagram 1.12 is an illustration of this process. Diagram 1.12

A 2sp2 hybridized atomic orbital of carbon linearly combines with a 1s atomic orbital of hydrogen to form a C-H sigma bond. Four of these C-H bonds are formed in the ethylene molecule. Each C-H single bond is described by Diagram 1.13.Diagram 1.13

The third 2sp2 hybridized atomic orbital linearly combines with the 2sp2 hybridized atomic orbital of the second carbon to form a 2sp2 + 2sp2 σ bond: The remaining electron in the 2p orbital of carbon linearly combines with the electron in the 2p orbital of the second carbon to form a π bond. These two types of linear combinations of atomic orbitals comprise the C-C double bond. Diagram 1.14 illustrates these linear combinations of atomic orbitals that form the molecular orbitals of the C-C double bond.Diagram 1.14

The overall pictorial representation of the ethylene molecule is illustrated in Diagram 1.15.Diagram 1.15

The pi bond can be described as the vertical overlap of two p atomic orbitals and the sigma bond can be described as the horizontal overlap of two 2sp2 hybridized atomic orbitals. Together, the pi and sigma bonds form the C-C double bond. Each C-H single bond can be described as the horizontal overlap of a 1s atomic orbital of hydrogen with a 2sp2 atomic orbital of carbon.The H-C-H bond angle is 120o and the H-C-C bond angle is 120o.

Multiple bond (the triple bond)

The simplest molecule with a triple bond is acetylene (ethene), C2H2. Following the octet rule, the structure for C2H2 would be:

A sigma, σ, and two pi, π, bonds comprises the triple bond.

In order to understand the nature of the triple bond, let’s revisit the electron configuration of each of the two carbons that comprise the acetylene molecule. The electron configuration of carbon is 1s22s22p2. In forming the triple bond, an electron in the 2s orbital of carbon is promoted to the 2p orbital of carbon. Diagram 1.16 is an illustration of this process.Diagram 1.16A 2p atomic orbital of carbon containing one electron linearly combines with the 2s atomic orbital to form two degenerate 2sp hybridized atomic orbitals.Diagram 1.17 illustrates this process.Diagram 1.17

A 2sp hybridized atomic orbital of carbon linearly combines with a 1s atomic orbital of hydrogen to form a C-H sigma bond. Two of these C-H bonds are formed in the acetylene molecule. Each C-H single bond is described by Diagram 1.18. Diagram 1.18

The second 2sp hybridized atomic orbital linearly combines with the 2sp hybridized atomic orbital of the second carbon to form a (2sp + 2sp) σ bond: The remaining two 2p orbitals of carbon linearly combines with the remaining two 2p orbitals of the second carbon to form two π bonds. Diagram 1.19 is an illustration of this process.Diagram 1.19

The overall pictorial representation of the acetylene molecule is illustrated in Diagram 1.20.Diagram 1.20

Two pi bonds are formed from the vertical overlap of four p orbitals and the sigma bond is the horizontal overlap of two 2sp hybridized atomic orbitals. of one carbon atom with another carbon atom. Together, the two pi bonds and the sigma bond comprise the C-C triple bond. Each C-H single bond can be described as the horizontal overlap of a 1s atomic orbital of hydrogen with a 2sp atomic orbital of carbon. The H-C-C bond angle is 180o.

Intramolecular forces

The two types of intramolecular forces encountered in organic systems are repulsive forces and attractive forces. Repulsive forces exist when electrons remain as far apart as possible because they have the same charge or the same spin. Attractive forces exist when the atomic nuclei attract electrons due to their opposite charges.

Bond Polarity

Bond polarity is a property of the covalent bond where the two nuclei don’t share the covalently bonded electrons equally. The electron cloud is denser about one atom more than the other; therefore, one end of the bond is relatively negative and the other is positive. Polarity is generally indicated in the following manner:

The properties of polarity

Bond polarity occurs when there is a difference in electronegativity between the atoms. The greater the difference in electronegativity, the more polar the bond. Fluorine is the most electronegative atom.

Electronegativity is a measure of the ability of an atom in a molecule to draw bonding electrons to itself.

Linus Pauling suggested electronegativity values from bond energies. However, to make this work, he had to arbitrarily assign an electronegativity value to the most electronegative element on the periodic chart, the fluorine, F, atom. Pauling assigned a value of approximately 4.0 to F, and the electronegativity values of the other elements were calculated using 4.0 as the electronegativity of fluorine.

The symbol for electronegativity is χ, and χ for an element can be calculated using the following equation:

Where χA is a known value

Example 1.1

Example 1.2

As indicated previously, electronegativity values are used to determine the extent of polarity in a bond. For example, The electronegativity of hydrogen is 2.1 and the electronegativity of chlorine is 3.0; therefore in HCl, the electronegativity difference would be 3.0-2.1 = 0.9. The electronegativity of fluorine is 4.0; therefore in HF, the electronegativity difference would be 4.0-2.1 = 1.9. Consequently, HF would have a more polar covalent bond than HCl.

Figure 1.12 the electronegativity of a variety of elementsFigure 1.12: Pauling Electronegativity of elementsFigure 1.12 was taken from the following Website: http://www.britannica.com/EBchecked/topic-art/108614/63/Each-element-has-an-electronegativity-value-which-is-a-measure#tab=active~checked%2Citems~checked

Polarity of Molecules

A molecule is polar when its center of negative charge does not coincide with its center of positive charge. Polar bonds in the molecule exhibit dipoles when two equal and opposite charges are separated in space. Under these circumstances the molecule exhibits a dipole moment, and the dipole is represented as

The arrow points from positive end of the molecule to the negative end of the molecule. The dipole moment of the molecule is represented by μ, where

The dipole moment depends upon the polarity of its individual bonds and the direction of the bonds ; however, calculating the dipole moment becomes more complicated when more than two atoms are involved. Carbon tetrachloride, compound III, is an example of a molecule that exhibits a zero dipole moment, and chloroform, compound IV, is a molecule that has a dipole moment. Remember that the dipole moment of a molecule is the vector sum of the bonds dipole moments (this is not easily calculated):

The net dipole moment of ammonia, compound V, is 1.46 D

NF3 compound VI, has a smaller dipole moment than NH3 because the differences in the directions of the bond dipoles:

Useful physical parameters of organic molecules include the dipole moment, the melting point, the boiling point, the solubility, and the density. These physical parameters provide information about organic compounds that may be helpful in identifying their structures.

The physical parameters can be used to determine methods of separating organic compounds from mixtures. Sometimes the isolation and purification methods are more complicated than the method of synthesis.

Resonance

Resonance is the phenomenon in which the actual structure is a hybrid between two or more structures.

Ozone exhibits the phenomenon of resonance.Ozone has the following two equivalent structures:

Resonance

Resonance is the phenomenon in which the actual structure is a hybrid between two or more structures.

Ozone has the following two equivalent structures:

The actual structure is not VI or VII, but a hybrid between VI and VII similar to VIII

Resonance in ozone can be represented by the following, where the double headed arrow represents resonance:

The Lewis structure for benzene is:

Acid and Bases

An Arrhenius acid is a substance that ionizes in aqueous solution to produce hydronium ions.

For example,

A Brønsted-Lowry acid is a substance that is a proton donor and a Brønsted-Lowry base is a proton acceptor.

For example,

The strength of the acid depends upon its tendency to give up protons.